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Author: Subject: method to check PH of methanol
ptlmayank
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[*] posted on 17-11-2010 at 03:39
method to check PH of methanol


what is the method to check the pH of methanol
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[*] posted on 17-11-2010 at 12:59


A pH meter.
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[*] posted on 18-11-2010 at 11:24


Why do you want to check the pH of methanol?
Just for a start, pure methanol does not have a pH of 7.

True pH measurements in non aqueous media are an absolute pain, using a pH meter is more likely to damage the pH probe than to give you the right answer..
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[*] posted on 19-11-2010 at 09:33


If you know the concentration of the MeOH in M and can find the Ka online then you can calculate it.
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[*] posted on 19-11-2010 at 14:45


How precise do you need the pH? Neutral litmus paper will give you a quick 'is it acidic or is it basic' heads up.



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[*] posted on 19-11-2010 at 16:37


Quantitatively mix in some water, and titrate the stuff. This will give you "titrable acidity" rather than "pH" (which is only defined in water, anyway). [H+] can be back-calculated.

Keep in mind that, in pure MeOH, the pKa of the hydroxyl proton which will be the limiter, e.g. a strong enough base is required to make the alkoxide or a strong enough acid is required to protonate the hydroxyl. This may yield a base, for example NaOMe, which is far stronger base than the corresponding NaOH created by quenching the mixture in water.

This is however, routine protocol for catalyst prep (and determination of free fatty acids in used food-oils) in the manufacture of biodiesel, etc.

Cheers,

O3




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[*] posted on 20-11-2010 at 02:54


The pKa of methanol is quite close to that of water, so for diluted aqueous solutions of methanol you can safely assume the pH is still very close to neutral (pH 7). With very accurate measurements you should see some difference, perhaps on at the second decimal point, but then even plain distilled water hardly gives an exact pH 7.00 measurement as any little contaminant can very easily alter that. For higher concentrations of methanol you can not truly speak of pH any more, because the pH is defined for aqueous solutions only (for example, a "70% aqueous solution of methanol" is not really aqueous any more).

Quote: Originally posted by Ozone  
Keep in mind that, in pure MeOH, the pKa of the hydroxyl proton which will be the limiter, e.g. a strong enough base is required to make the alkoxide or a strong enough acid is required to protonate the hydroxyl. This may yield a base, for example NaOMe, which is far stronger base than the corresponding NaOH created by quenching the mixture in water.

Contrary to popular belief, this is not true. In water, the methoxide ion is a very slightly weaker base than hydroxide (though admittedly the difference is small - hydroxide is only 1.6-times more basic than methoxide). However, in methanol/ethanol the hydroxide ion is substantially more basic than the methoxide/ethoxide (there is a thread in the General section where a paper on exact measurements was posted). Also in general, in most organic solvents, due to the obvious solvation and polarizability reasons, the alkoxide ions of primary alcohols are always less basic than the hydroxide ion.

[Edited on 20/11/2010 by Nicodem]




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[*] posted on 20-11-2010 at 11:18



' "pH" (which is only defined in water, anyway)'
"For higher concentrations of methanol you can not truly speak of pH any more, because the pH is defined for aqueous solutions only "


Why do people think that the log of the reciprocal of the hydrogen ion activity is only defined in aqueous solutions?

It's a pig to measure in, for example, methanol, but it's not impossible.
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[*] posted on 20-11-2010 at 13:29


Quote: Originally posted by unionised  
using a pH meter is more likely to damage the pH probe than to give you the right answer..


why is this?

methanol isn't particularly corrosive
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[*] posted on 20-11-2010 at 14:10


This discussion makes it clear that "simple"concepts such as pH are not "clear" at all. I have attached a nice paper which, at least, clearly defines the problem.

[Nicodem]The pKa of methanol is quite close to that of water, so for diluted aqueous solutions of methanol you can safely assume the pH is still very close to neutral (pH 7). With very accurate measurements you should see some difference, perhaps on at the second decimal point, but then even plain distilled water hardly gives an exact pH 7.00 measurement as any little contaminant can very easily alter that.

Correct, distilled/deionized water will abscond with ions from wherever, e.g. particles in the air, etc. The lack of ions canalso "confuse" a standard pH probe. Mine (18 MO) reads acidic at
6.8, or so, but has essentially zero buffering capacity (a topic for a separate discussion).

[Nicodem]For higher concentrations of methanol you can not truly speak of pH any more, because the pH is defined for aqueous solutions only (for example, a "70% aqueous solution of methanol"
is not really aqueous any more).

I did point out that pH is defined in aqueous solution. This is why I recommended that it be diluted, eg. 1:1. I then recommended titration because I don't like the use of "pH" with respect to non-aqueous solvents; "titrable acidity" is a normalizing metric. However, in methanol, in particular (as you point out, see discussion below), the pH probe will likely suffice. This is not so true, though, in other solvents that may destabilize anions leading to much greater
basicity, e.g. DMSO.

I should have taken greater care to point out that, from

http://en.wikipedia.org/wiki/Acid_dissociation_constant :

"When a compound has limited solubility in water it is common practice (in the pharmaceutical industry, for example) to determine pKa values in a solvent mixture such as water/dioxane or water/methanol, in which the compound is more soluble.[27] In the example shown at the right, the pKa value rises steeply with increasing percentage of dioxane as the dielectric constant of the
mixture is decreasing.

A pKa value obtained in a mixed solvent cannot be used directly for aqueous solutions. The reason for this is that when the solvent is in its standard state its activity is defined as one. For example, the standard state of water:dioxane 9:1 is precisely that solvent mixture, with no added solutes. To obtain the pKa value for use with aqueous solutions it has to be extrapolated to zero co-solvent concentration from values obtained from various co-solvent mixtures."

Which was the extrapolated correction I was alluding to earlier. At least the hydrolysis of methoxide, if it exists in that mixture, will normalize (mole-wise) to hydroxide in water, which we can measure with some confidence, despite the required correction.

[Nicodem]Contrary to popular belief, this is not true. In water, the methoxide ion is a very slightly weaker base than hydroxide (though admittedly the difference is small - hydroxide is only 1.6-times more basic than methoxide). However, in methanol/ethanol the hydroxide ion is substantially more basic than the methoxide/ethoxide (there is a thread in the General section where a paper on exact measurements was posted). Also in general, in most organic solvents, due to the obvious solvation and polarizability reasons, the alkoxide ions of primary alcohols are always less basic than the hydroxide ion.

Oops. Yes, this is true (my reference cites 1.8 times and pKa water = 15.7, MeOH = 15.5).

The definition of a superbase(http://en.wikipedia.org/wiki/Superbase ) can get you into trouble
and even advanced organic chemistry neglects to mention hydroxide in non-aqueous environments, it's simply, "strong base", of which methoxide is the simplest case.

Methoxide cannot exist in water (that I know of) because it hydrolizes to yield NaOH. Thus, the practical comparison of basicity of MeO- to -OH in water is moot.

So, I'll agree that if hydroxide can exist, in the dissociated state, in Methanol [which it appears to do, albeit in small concentrations (I can't find a decent reference for this), which appear relative to a very small amount of contaminating water], it would be a greater base than methoxide.

I'll posit, then, that the basicity (or acidity) is limited by the concentration and activity of the active species in the media in which it is measured, so long as it can exist in measurable concentration.

At any rate, for methanol dilute (greater than 40 mol %--from attached paper) with water and use a probe or titrate. Correct [H+] for dilution and recalculate "pH".

Thanks,

O3

Attachment: Abrash 2001.pdf (84kB)
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[*] posted on 21-11-2010 at 10:13


Quote: Originally posted by spirocycle  
Quote: Originally posted by unionised  
using a pH meter is more likely to damage the pH probe than to give you the right answer..


why is this?

methanol isn't particularly corrosive


Well, since it almost certainly won't work it's more likely to damage it than to work.
The glass membrane that acts as a sensing medium for most pH probes is designed to be used wet with water. That's why you shouldn't let them dry out.
Putting it in methanol will dry it very quickly.

Also, all this discussion about dilute solutions of methanol is probably off-topic since the original question was about methanol, not water with a bit of methanol in it.
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